Magnesium sulphite is an important intermediate in many large-scale chemical processes, such as the desulphurization of gases and air pollutants by the magnesite process. However, magnesium sulphite is in many cases considerably contaminated by solid impurities from the starting magnesite and other impurities from flue gases. Industrial applications require pure magnesium sulphite, and purification methods have been developed to meet this need.
The removal of impurities from easily soluble crystalline substances on the basis of solubility differences occurring at different temperatures is known. J. W. Mullin, Crystallization, CRC Press. These processes relate to the separation and subsequent crystallization of dissolved substances in solutions having concentrations that are as close as possible to the equilibrium determined by the temperature and pressure to which the solution is exposed. A disadvantage of this method is that it requires a considerable amount of time for the dissolving and crystallization steps, in order to obtain a sufficient yield.
Magnesium sulphite is highly insoluble; only 0.7% by weight readily dissolves. The magnesite impurities contained in raw magnesium sulphite are also highly insoluble. The equilibrium concentrations of these materials in solution is thus very low. As a result, traditional separation methods based on differential solubility in water (as taught by Mullin) have not been successful, and are not used in the industry.
It is also known that the solubility of some substances (such as MgSO.sub.3) is anomalous under certain temperature-dependent conditions, and that transient supersaturated metastable solutions can arise when these conditions are met. See Czechoslovak Author's Certificate No. 209,952. However, such metastable solutions have been unpredictable, very short lived, difficult to control, and they rapidly revert to equilibrium solutions by crystallization of the excess. This phenomenon has not been efficient when used in a commercial process because of the serious recrystallization problem. The known process requires the rapid dissolution of magnesium sulphite crystals in water, and the rapid separation of impurities from the resulting warm solution. The preparation of this solution on an industrial scale has met with many difficulties. The required supersaturated solution could not be induced at all upon dissolution in a stirred heated charging reactor. Instead, MgSO.sub.3.6H.sub.2 O was immediately recrystallized into MgSO.sub.3.3H.sub.2 O.
Dissolution in a through-flow tubular reactor produced a supersaturated solution, but only for a very short time after start-up. The concentration of the solution then decreased, and the device had to be taken out of operation after several hours because the piping became encrusted with MgSO.sub.3.3H.sub.2 O crystals. A major disadvantage of the known method resides in the undesirable conversion of hexahydrate into trihydrate, with the resulting accumulation of trihydrate in the pipes during heating. This causes interruption of the process for periodic and laborious cleaning of the pipes.
The known process is successful only if the extraction and separation of impurities can be achieved rapidly, within the short life of the metastable solution. In addition, the known process is disadvantageous because it requires that the raw hexahydrate be free of magnesium sulphite trihydrate, since the presence of trihydrate produces a seeding effect which prohibits the extraction of hexahydrate and produces a poorly soluble trihydrate in industrial plants at temperatures of from 60.degree. to 120.degree. C.
Other apparently promising methods of obtaining a useful metastable solution have not been successful. Indirect heating in a continuous heat exchanger results in rapid fouling and clogging of the exchanger with incrustations of trihydrate. Precipitate also accumulates in a batch reactor, or in a batch reactor used in combination with a heat exchanger. In order to obtain the most efficient use of the metastable solution, it is necessary to minimize (and if possible completely avoid) the gratuitous conversion of hexahydrate into trihydrate. This objective is particularly difficult in the known processes, because the temperatures they disclose encourage the formation of an undesired stable trihydrate solid in equilibrium with the desired liquid solution. Trihydrate precipitation and its later removal have therefore become recognized as necessary though undesirable features of the art processes.